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Oxidation and reduction

Introducing oxidation states, oxidation, and reduction. Some tips for remembering oxidation and reduction.  Created by Sal Khan.

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  • old spice man green style avatar for user aloukik mishra
    for a redox reaction to occur oxidation and reduction must occur together or could one
    would be considered a redox reaction
    (67 votes)
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    • leafers tree style avatar for user Thane McGarry
      Both must occur together for it to be a redox reaction, but if only one occurred it would be called an "oxidation half-reaction" or a "reduction half-reaction." Finding these "half reactions" are often essential to solve a redox reaction problem, but they are NOT redox reactions themselves. I hope that helps!
      (18 votes)
  • mr pink red style avatar for user Rainman
    why H-F dont make ionic bond? As H has one valence electron and F has 7 ??
    (21 votes)
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    • piceratops ultimate style avatar for user Just Keith
      Hydrogen makes very few ionic compounds.

      If you want to think in terms of the octet rule, H would have no electrons whatsoever if it gave its electron to F. So, for H to achieve its full two electrons (remember that H seeks 2, not 8, electrons) then the only way to do that with F is by sharing the electron. Thus a polar covalent bond results in F and H both having what they want.

      And, in fact, all of the hydrogen halides in their pure form have polar covalent or covalent bonds. None is ionic. However, when dissolved in water to form acids, the chemistry is a bit different. HF is a weak acid and HCl, HBr, and HI are strong acids.

      So, it is rather usual for hydrogen to be involved in an ionic bond. There are a few such compounds (some of the hydrides are ionic), but it is unusual.

      I think I should note that some websites get this wrong. There is an article on eHow that really messes this up. Some mistakenly think that just because some compounds of hydrogen easily dissociate in water (the strong acids) then they must be ionic. But, this is not the case. There is a chemical interact between the hydrogen halides and water that allows all but HF to fully dissociate, but the molecules in their pure form are not ionic.

      By the way, hydrogen is not the only element that unexpectedly doesn't make many ionic compounds. In particular, despite being in Group 2, NONE of beryllium's compounds are ionic. Even BeF₂ is not ionic. I've seen several websites the get this wrong.

      The important thing to learn here is the the rules of thumb about what will or will not be an ionic compound do not apply to every situation. Even the method of electronegativity difference, though often quite helpful, doesn't always work.
      (45 votes)
  • leaf yellow style avatar for user Abdullah Hussein
    So basically, what I understood is:

    1. loss of electrons,
    2. increase in oxidation state.

    1. gain of electrons,
    2. decrease in oxidation state.

    Right?! :)
    (11 votes)
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  • male robot donald style avatar for user Mit Vasani
    isnt the bond in water ionic? as 2h+ and o2-?
    (7 votes)
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  • blobby green style avatar for user Jennifer D
    What are the differences between formal charge (which we write the sign after number, 2+) and oxidation number (which we write the sign before number, +2)?
    (6 votes)
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    • spunky sam blue style avatar for user Ernest Zinck
      In formal charge, the shared electrons are counted as if they were shared equally. In H-O-H, each atom gets one electron from the shared bond. H normally has one valence electron. It has neither gained nor lost electrons, so its formal charge = 0. The O atom gets one electron from each O-H bond, plus its four lone-pair electrons, for a total of six electrons. It normally has six valence electrons. It has neither gained nor lost electrons, so its formal charge = 0.

      In oxidation number, the shared electrons are counted as if they belong entirely to the more electronegative atom. So the O atom gets all the shared electrons and H gets none. H normally has one valence electron; now it has none. So its oxidation number is +1. The O atom gets two electrons from each O-H bond, plus its four lone-pair electrons, for a total of eight electrons. It normally has six valence electrons, so its oxidation number is -2.
      (13 votes)
  • blobby green style avatar for user Manni Desruisseaux
    When you say that hydrogen is being oxidized by oxygen, does that mean that oxygen is the oxidation agent, and hydrogen would be the reduction agent? Am I getting this right?
    (7 votes)
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  • leaf red style avatar for user http://facebookid.khanacademy.org/100004343314752
    how can we know charge of an ion
    (6 votes)
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    • piceratops ultimate style avatar for user Just Keith
      You determine the charge of an ion by adding the total number of protons and subtracting the total number of electrons that are actually present, then you will get the charge of the ion. However, you would be well advised to simply memorize the charges on all of the more commonly encountered ions.
      (7 votes)
  • blobby green style avatar for user Erica Mitchell
    at , what is an example of a halogen?
    (2 votes)
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  • ohnoes default style avatar for user Marcus Fong
    Can the oxidation state be also referred to as the oxidation number?
    (4 votes)
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  • leaf red style avatar for user Peter Theobald
    From around , I have realised that for something to be reduced something has to be oxidized and vice-versa .

    Thought ?
    (2 votes)
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Video transcript

Let's think a little bit about the molecule sodium chloride. If we look at the periodic table right over here, we see that sodium is a Group 1 element. It's an alkali metal. It has one valence electron. It's also not too electronegative. It's sitting here on the left-hand side of the periodic table. We know the general trend for electronegativity is that it increases as we go to the top right. So these elements over here generally like to hog electrons. These elements over here generally like to give away electrons. These are electropositive, these are electronegative. So sodium right over here, with its one valence electron, it's a very good candidate for giving away an electron. On the other side of the periodic table, we have chlorine right over here. It's a Group 7 element. It's a halogen. It would love nothing more than to gain an electron so that it can get to the magic number of eight valence electrons. It is very electronegative. So what you can imagine would happen is that if these two things were to interact? The chlorine would nab sodium's extra electron over here, and it would become the chloride anion. And then sodium would become the sodium cation. Cation just is a positive ion. Anion is just a negative ion. And so this one right over here is now positive. Sodium is now positive. Chloride is a negative ion. And so they're going to be attracted to each other, and they're going to form this ionic bond. So this right over here is an ionic bond. Now let's think about something that's not ionic, where the electrons are not being fully nabbed from one atom to another, but they're being shared. And one of the most famous examples of that is water. So we know water is H2O. Each water molecule is one oxygen bonded to two hydrogens. And these two bonds are covalent bonds. In each of these bonds, you have a pair of electrons that are being shared by both the hydrogen and the oxygen. But we also know that this isn't a completely equal sharing of the electrons. We look on the periodic table here, oxygen is far more electronegative than hydrogen is. And so because of that, the electrons in these two bonds are going to spend more time around oxygen than they are going to spend around the hydrogens. And we've seen this before. This would give the oxygen end of the water molecule a partially negative charge. That's the lowercase Greek letter delta. We use that for the notation of partially negative. And on the hydrogen ends of the molecule, we're going to have a partially positive charge. Now, this is the reality. But as we'll see later on in future videos, it's sometimes inconvenient to have this partial messiness. And so what I'm going to do right now is introduce you to what is fundamentally just an intellectual tool. It's just a convention that chemists have invented that allow us to get our heads around a lot of reactions and allow us to think about how is a reaction likely to occur. And that intellectual tool is the idea of oxidation states. What the oxidation state is, even if you're in a situation where you have covalent bond, you say, well, look, I understand. Those are partial charges. These are covalent bonds. The electrons are being shared. But I don't like this partial stuff. I want to just assume hypothetically, what if these were ionic bonds? And you say, well, if these had to be ionic bonds, then the oxygen would nab the electrons from these pairs. And so the oxygen would have a fully negative charge, a negative 2 charge. And the hydrogens would have a fully positive charge each. And so, if we were to write down the oxidation states for the atoms in the water molecule-- let's write that down, so H2O-- we would say that oxygen has an oxidation state of negative 2, and each hydrogen atom has an oxidation state of plus 1. And notice, the whole molecule is neutral, and these things cancel out with each other. Positive 1, positive 1, that gets you to positive 2. Then you have negative 2. They cancel out. Now, the one thing, I keep saying this is negative 2, but I wrote the negative after it. If I wanted to write positive 1 as an oxidation state, I would actually write it as 1 positive, although you can assume that if someone just writes the positive. And this is just the convention, to write the sign after the number when we are writing actually ionic charges or oxidation states, because an oxidation state is nothing but a hypothetical ionic charge. If you really had to-- if you were forced to assume these aren't covalent bonds, but these are ionic bonds. Once again, I want to stress. This is the reality. These are covalent bonds. These are partial charges, the oxidation state, intellectual tool, that's forcing us to pretend like these are ionic bonds. And you might say well, this kind of makes sense right over here. This involved oxygen in some way. That's why it's called oxidation states. And that's how I initially conceptualized it when I first learned about this. You say, well, look, each of these hydrogens lost an electron to oxygen. So it makes sense that we say that each hydrogen got oxidized, so hydrogen oxidized by oxygen. It makes sense that oxygen would oxidize something else. This got done to it. The charge was taken away by oxygen, so it got oxidized. Now, the other term on the other side of oxidized is reduced. And the word "reduced" really comes from the idea that oxygen's charge has been reduced. So we could say, O, or we could say oxygen has been reduced by the hydrogens. And so there there's a temptation here to say, well, OK, this must always involve oxygen in some kind, because it seems to begin with the same words. Well, that is not the case. Let's take, for example, if this is an aqueous solution, hydrofluoric acid right over here. You have a hydrogen covalently bonded to a fluorine. Now, just like we saw in water, fluorine is one of the most electronegative elements. It's going to hog the electrons in this covalent bond. So this is going to have a partial negative charge here. This is going to have a partial negative charge here. And this is going to have a partially positive charge. But if we were going to think about it in terms of oxidation states, we would say when push comes to shove, if this had to be at an ionic bond and not a covalent bond, what would be the charges on each of these atoms? We can say, well, in that case, hydrogen would lose an electron, and it would have a full positive charge. And fluorine would gain an electron and have a full negative charge. This is a hypothetical. Once again, the reality is they're partial. It's a covalent bond. But the hypothetical one is a full positive charge here and a full negative charge here. And so we would say that the oxidation state in this molecule for hydrogen is plus 1 and the oxidation state for fluorine in this molecule is negative 1. And we would say that hydrogen has been oxidized. And we could even say it's been oxidized by the fluorine. And we would say that the fluorine, because its hypothetical ionic charge has been reduced, we would say fluorine might be reduced. And you would say, wait, wait. Look, oh wow, we're using the word "oxidize" even though there is no oxygen to be seen here. And one way to think about it, if someone had told that you had been Bernie Madoffed, that doesn't necessarily mean that you interacted with Bernie Madoff. It means that someone did to you what Bernie Madoff would have done. Someone else-- someone took your money, told you they were going to invest it, and then put it into a Ponzi scheme. Even if that person is not Bernie Madoff, you could say that you have been Bernie Madoffed. So here, fluorine did to hydrogen what oxygen tends to do. It took an electron away. It oxidized the hydrogen. Now, that's how I tend to remember it. If something has been oxidized, it's losing an electron. And what I think about it is, well, that's what oxygen would have done to you. Oxygen is very electronegative. It tends to take electrons away from other atoms. Now, there are other mnemonics that you might see for remembering what oxidation and reduction actually represent. And I'll introduce those to you, just because they might be helpful, and they are introduced in a bunch of chemistry classes. One of the mnemonics is LEO the lion says GER. I'll write says in lowercase here, because it's really not relevant. But LEO the lion says GER. And this is to remember that losing an electron means you are being oxidized, or losing electrons is oxidation. And gaining electrons is reduction. So that's just a mnemonic. Another one that's often used is OIL RIG. And this, essentially-- oxidation is losing electrons, reduction is gaining electrons.